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Calculate buffer pH using pH = pKa + log([A⁻]/[HA]) or find the conjugate base to acid ratio for any target pH. Includes pKa values for common biological buffers.
Buffers resist pH changes when acid or base is added. Henderson-Hasselbalch relates buffer pH to pKa and the ratio of conjugate base to acid. Blood, cells, and enzyme reactions all require precise buffering.
| Buffer | pKa | Effective Range | Use |
|---|---|---|---|
| Acetic acid/Acetate | 4.75 | pH 3.8-5.8 | Food, biochemistry |
| Phosphate (pKa₂) | 7.20 | pH 6.2-8.2 | PBS, biology |
| HEPES | 7.48 | pH 6.8-8.2 | Most popular bio buffer |
| Tris-HCl | 8.07 | pH 7.0-9.0 | Most common lab buffer |
| Bicarbonate | 6.35 | pH 5.5-8.0 | Blood buffer |
Weak acid (HA) + conjugate base (A⁻). Add H⁺: neutralized by A⁻. Add OH⁻: neutralized by HA. Best at pH=pKa. Effective within pKa ± 1.
pH 7.35-7.45. Bicarbonate buffer: lungs adjust CO₂ (fast), kidneys adjust HCO₃⁻ (slow). pH 7.0 causes fatal arrhythmias.
137 mM NaCl, 2.7 mM KCl, 10 mM Na₂HPO₄, 1.8 mM KH₂PO₄, pH 7.4. Most common cell biology buffer.
Buffer solutions are indispensable in chemistry and biology. They maintain stable pH in reaction media, protect biological systems from pH swings during metabolism, preserve the integrity of analytical measurements, and enable precise control of chemical reactions. Understanding buffer chemistry — and especially the Henderson-Hasselbalch equation — is essential for any chemist or biologist working with aqueous systems.
Buffer capacity (β) measures how much acid or base a buffer can absorb before significant pH change. β = dn/dpH where n is moles of acid/base added per liter. Maximum buffer capacity occurs at pH = pKa (equal amounts acid and base). Buffer capacity formula: β = 2.303 × C × Ka[H⁺]/(Ka + [H⁺])². Where C is total buffer concentration. A 1 M buffer has 10× the capacity of a 0.1 M buffer at the same pH. This is why high-capacity buffers are used when large amounts of acid or base will be added.
Biological research uses synthetic buffers (Good buffers, developed by Norman Good in 1966) because many classical buffers interfere with biology. MES (pKa 6.15): cell cultures at acidic pH. PIPES (pKa 6.76): intermediate pH applications. HEPES (pKa 7.48): most widely used biological buffer, does not penetrate cell membranes, minimal metal binding. Tris (pKa 8.07): very common in molecular biology (DNA electrophoresis buffers, protein gels). Bicarbonate (pKa 6.35): physiological buffer in CO₂ incubators — 5% CO₂ atmosphere maintains pH 7.4 in bicarbonate-buffered media.
Drug stability and efficacy often depend critically on pH. Aspirin (pKa 3.5) is fully ionized (negatively charged) at blood pH 7.4, preventing it from re-crossing cell membranes — important for its pharmacokinetics. Insulin is most stable at pH 4 (acid formulations) but is injected and distributed at blood pH 7.4. Many drugs are formulated as their conjugate acid or base (hydrochloride or sodium salt) to improve solubility. Injectable drugs must be formulated within narrow pH ranges (pH 3-9) to avoid pain and tissue damage at injection sites.
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A buffer solution is a mixture of a weak acid and its conjugate base (or a weak base and its conjugate acid) that resists changes in pH when small amounts of strong acid or base are added. The Henderson-Hasselbalch equation pH = pKa + log([A⁻]/[HA]) quantitatively describes the relationship between buffer pH, the acid dissociation constant (pKa), and the relative concentrations of conjugate base [A⁻] and weak acid [HA]. This equation is among the most practically useful in all of chemistry and biochemistry.
Buffers are essential wherever stable pH is required: biological systems (blood, cell culture media, enzyme reactions), analytical chemistry (standardizing instruments, preserving sample pH), pharmaceutical formulations (drug stability depends on pH), food preservation, and industrial processes from electroplating to fermentation.
A buffer containing weak acid HA and conjugate base A⁻ resists pH changes by two complementary reactions. When strong acid (H⁺) is added: H⁺ + A⁻ → HA. The conjugate base consumes the added acid, preventing H⁺ accumulation and pH drop. When strong base (OH⁻) is added: OH⁻ + HA → A⁻ + H₂O. The weak acid neutralizes the added base, preventing OH⁻ accumulation and pH rise. The reservoir of both HA and A⁻ provides buffering capacity. Buffer capacity (β) is highest at pH = pKa (equal concentrations of acid and base) and increases with total buffer concentration.
Starting from the Ka expression for weak acid dissociation HA ⇌ H⁺ + A⁻: Ka = [H⁺][A⁻]/[HA]. Rearranging: [H⁺] = Ka × [HA]/[A⁻]. Taking -log of both sides: -log[H⁺] = -log(Ka) - log([HA]/[A⁻]). Therefore: pH = pKa + log([A⁻]/[HA]). When [A⁻] = [HA] (equal concentrations): log(1) = 0, so pH = pKa. This is why buffers work best at pH = pKa — they contain equal amounts of acid and base, maximizing buffering capacity in both directions. Effective range: pKa ± 1 (within this range, [A⁻]/[HA] ratio stays between 0.1 and 10).
To prepare 1 L of 50 mM phosphate buffer pH 7.4 using Na₂HPO₄ (pKa₂ = 7.20) and NaH₂PO₄: Apply Henderson-Hasselbalch: 7.4 = 7.20 + log([HPO₄²⁻]/[H₂PO₄⁻]). log ratio = 0.20. Ratio = 10^0.20 = 1.585. Total = 1.585 + 1 = 2.585 parts. HPO₄²⁻ = 50 mM × 1.585/2.585 = 30.67 mM: dissolve 30.67 mmol × 141.96 g/mol = 4.352 g Na₂HPO₄ in 800 mL water. H₂PO₄⁻ = 50 - 30.67 = 19.33 mM: add 19.33 mmol × 119.98 g/mol = 2.319 g NaH₂PO₄. Dilute to 1 L. Verify pH with calibrated electrode. Adjust with 1M NaOH or HCl if needed. Always verify pH with electrode — the Henderson-Hasselbalch equation gives a starting point, but temperature and activity effects may require small adjustments.
Blood pH is maintained at 7.35-7.45 by three buffer systems: (1) Bicarbonate/CO₂ buffer (most important, fastest): CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻. pKa = 6.10. Normal blood: [HCO₃⁻] = 24 mEq/L, [H₂CO₃] = 1.2 mEq/L (determined by dissolved CO₂). pH = 6.10 + log(24/1.2) = 6.10 + 1.30 = 7.40. Lungs control [CO₂] (fast, seconds-to-minutes response). Kidneys control [HCO₃⁻] (slow, hours-to-days). (2) Hemoglobin buffer: deoxyhemoglobin is a better base than oxyhemoglobin — tissue CO₂ is buffered as hemoglobin releases O₂. (3) Phosphate buffer: intracellular pH buffering (pKa₂ = 7.20, ideal for intracellular pH ~7.2).
Buffer capacity β = d[base added]/dpH (differential moles of base per pH unit change). For the Henderson-Hasselbalch buffer: β = 2.303 × C × Ka[H⁺]/(Ka + [H⁺])², where C = total buffer concentration. Maximum capacity at pH = pKa: β_max = 2.303 × C/4 = 0.576 × C. A 0.1 M acetate buffer at pH 4.74 has β_max = 0.576 × 0.1 = 0.0576 mol/(L·pH unit). Adding 0.05 mol/L strong acid changes pH by 0.05/0.0576 ≈ 0.87 pH units. Doubling buffer concentration doubles capacity. This is why concentrated buffers (0.1-0.5 M) are used when large amounts of acid or base will be added (cell culture maintaining pH while cells produce acidic metabolites).